M. L. MILLERAND C. L. SHERIDAN
184
linear phosphates AMP, ADP and ATP, the values for K’ for AQP were predicted using the equation K‘ = nu n2b and nac where n is the number of phosphate residues. The values of a, b and c were calculated from the apparent stability constants for AMP, ADP and ATP, and the values of K’ for AQP listed as “predicted” in Table I11 were obtained by letting n = 4. The uncertainty indicated is-that caused by K‘ being either too high or too low for AMP, ADP and ATP by the indicated uncertainties. The fact that in all cases the experimental values fall in the expected range supports the linear structure for the phosphate chain of AQP. A compafison of t,he pKa’ values also offers evidence for the linear structure. On the basis of the quite regular increases in pKa’ values in the series AMP, ADP and ATP, it was predicted that the values of PKa‘ for AQP would be 7.26 in (nC3H7)4NBr and 7.14 in (C2H&NBr. Since appre-
+
Vol. 60
ciable concentrations of complexes are formed in the case of K+, Na+ and Li+, this simple extrapolation procedure is inadequate, and it is necessary to consider, in addition, the expected increase in the stability constants of the complexes formed. The values for pKa’ for AQP in 0.2 M KC1, NaCl and LiCl predicted on the basis of the predicted pKa’ value in (n-CaH7)4NBrand the predicted stability constants in Table I1 are 6.57, 6.43 and 5.92, respectively. Thus the pKa’ values for AQP obtained experimentally agree within the experimental uncertainties with the values expected for a linear phosphate. Acknowledgments.-The authors are indebted to Dr. Robert M. Bock for assistance in setting up the recording p H meter and designing the automatic buret. This research was supported by the Research Committee of the Graduate School of the University of Wisconsin from funds supplied by the Wisconsin Alumni Research Foiindation.
CONCENTRATED SALT SOLUTIONS. I. ACTIVITY COEFFICIENTS OF SODIUM THIOCYANATE, SODIUM IODIDE AND SODIUM PERCHLORATE AT 25’ BY M. L. MILLER AND C. L. SHERIDAN Cmlribution from the Stam.ford Laboratories, Research Division,American Cyanamid Company, Stamford, Cmn. Received April 3.9,1966
Isopiestic vapor pressure measurements at 25” have been extended to near saturation for NaCNS, NaI and NaCIO,. With these data, values of the osmotic coefficients and activity coefficients up to near satura.tion have been computed.
I. Introduction Although the behavior of aqueous solutions of 1-1 electrolytes has been intensively investigated in the dilute and moderately concentrated range (up to 5 m) there have been fewer measurements in very concentrated solution (above 5 m). This is in part due to the limited number of salts which form such concentrated solutions and in part to a lack of interest in the field of very concentrated solutions. The isopiestic vapor pressure method developed by Robinson and Sinclairl and perfected by Robinson2is admirably suited for the measurement of activity coefficients in very concentrated solutions. By this method, measurements have been made on LiCP up to 15 m, LiBr4to 20 m and NH4NOa6to 26 m. Three other salts whose solubility extends up into the very concentrated range are NaSCN, NaI and NaC104. Measurements on these salts have, heretofore, been carried only as far as KC1 or NaCl could be used as a reference salt, i.e., t o 4, 5 or 6.5 m.2.6.7
As a part of a study of concentrated salt solutions (1) R. A. Robinson and D. A. Sinclair, J . A m . Chem. SOC.,66, 1830 (1934). (2) R. A. Robinson, {bid., 67, 1161, 1165 (1935). (3) R. A. Robinson, Trans. Faraday SOC.,41, 756 (1945). (4) R. A. Robinson and H. J. MoCoach, J . A m . Chsm. Soc., 69, 2244 (1947). (5) B. F. Wishaw and R. H. Stokes. Trans. Faradav SOC.,48, 27 (1953). (13) R. A. Robinson, J . A m . Chem. Soc., 6% 3131 (1940). (7) J. H. Jonea, THrs JOURNAL, 61, 516 (1947).
in this Laboratory, we have extended the vapor pressure measurements on these three salts up to near saturation. For this purpose, we have used sulfuric acid as a reference electrolyte and the vapor pressure measurements of Shankman and Gordons as standard. Other properties of the concentrated solutions of these salts which we have measured and which will be reported in forthcoming papers of this series are : viscosity, density, electrical conductivity and diffusion. 11. Experimental Method of Measurement.-The method of measurement was a somewhat modified version of that of Robinson and Sinclair.1 A 10” glass desiccator was used for the tests. This contained a */e‘‘ stainless desiccator plate into which had been drilled 4 circular depressions 1/4” deep t o hold the solution cups. The solution cups, two each for salt solution and sulfuric acid, were low-form glass weighin bott!es 2.5 inches wide and 1” high (outside closure). 8olutions were weighed into these bottles which were placed in the depressions on the desiccator plate. The covers were placed beside the bottles. This could be done very rapidly. The desiccator was sealed with silicone lubricant. The whole apparatus was placed on a rofking platform in a constant temperature room a t 25 1 . Vacuum waa riot used because it was found not to shorten the equilibrating time which ranged from 3 days to 2 weeks. Approximately half of the measurements were made with distillation into the sulfuric.acid and half with distillation out of it. After equilibrium had been attained, all solutions lost weight on successive weighings. The relatively long time required to reach equilibrium resulted from the use of glass (8) 8. Shankinan and A. R. Gordon, J . Am. Chem. SOC.,61, 2370 (1939).
ACTIVITY COEFFICIENTS OF CONCENTRATED SALTSOLUTIONS
Feb., 1956
185
instead of metal cups.a The long equilibration time was TABLE I preferred to the risk of corrosion from sulfuric acid. CONCENTRATIONS OF ISOTONIC SOLUTIONS Preparation of Solutions.-In working with concenMolality of Molality trated solutions of salt, it was found convenient to establish sulfuric acid, of salt, a precise curve for density versus concentration and to use moles 1000 moles/1000 this in determining the concentration of solutions3 subseg. Ha0 g. $0 Salt quently made up. 3.06 4.40 Sodium thiocyanate Sodium Thiocyanate.-Three different samples of so4.51 6.40 dium thiocyanate were used, namely, (1) J. T. Baker analyzed; (2) J. T. Baker analyzed, recrystallized from eth4.80 7.22 anol; (3) Mallinkrodt C.P. Density values for samples 5.77 8.36 2' and 3 were found t,o fall (within 0.64%) on the smooth 6.86 10.30 curve for sample 1. Therefore, J. T. Baker analyzed was 7.80 12.44 used without further purification for the vapor pressure work. Densities were determined a t 29.87 f 0.02' by 8.50 14.11 conventional pycnometric techniques. The temperature 9.12 15.20 of the bath was measured on a platinum resistance thermome17.86 10.15 ter calibrated a t the National Bureau of Standards. The 19.01 10.85 concentrations of samples for density work were determined by both gravimetric and volumetric methods. 2.32 Sodium iodide 3 . 9 4 Density values are These analyses agreed to &O.l%. 5.16 7.10 given in Paper I1 of this series. 6.54 9.19 Sodium Iodide.-Sodium iodide solutions were made up without additional purification from reagent salt supplied 6.89 9.29 by City Chemical Corporation. Concentrations were de7.96 11.90 termined gravimetrically, and also by measuring the den8.52 12.72 sity and referring to the density data in International Critical Tables (at 30" corrected by a small factor to 29.87'). The 15.14 Sodium perchlorate 7.65 agreement between the concentrations as measured by these 8.11 16.25 two methods was taken as evidence that the Ralt was of 5.81 10.85 sufficient purity. This procedure was considered safer than 6.14 3.88 attempting to dry a salt as labile as sodium iodide. TABLE I1 ACTIVITY A N D ONE MINUSOSMOTIC COEFFICIENTS Moles/1000 g. HaO. m
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
NaSCN Y
[0.712] .751 .82 .92 1.03 1.16 1.31 1.40 1.69 1.90 2.03 2.30 2.46 2.63 2.8 2.9 3.1 3.2
NaI
1 - 9
+0.031 -0.029 .003 .154 ,219 .285 .355 .422 .484 .538 ,578 .602 .619 .627 .633 .638 .642 .644
[0.736] .820 1.05 1.25 1.72 2.23 2.88 3.68 4.69 5.94 7.4 9.0
Sodium Perchlorate.-Sodium perchlorate solutions were made up by weight from Hodium perchlorate ("anhydrous") supplied by City Chemical Corporation. The salt was dried to constant weight in an oven a t 110'. No further purification was attempted. Examination of the salt bv ultraviolet emimion spectroscopys showed the presence of not over 0.1% of any one of 38 metal or metalloid impurities looked for.
111. Results
Table I gives the concentrations of the isotonic solutions of sodium thiocyanate, sodium iodide and sodium perchlorate as measured against sulfuric acid. In calculating activity coefficients from these data, reference values a t 1 m were taken from the compilation of Robinson and Stokes'O for sodium thiocyanate and sodium iodide. For sodium (9) Kindly carried out by Mr. W. L. Dutton of this Laboratory. (10) R. A. Robinson and R. H. Stokes, Trans. Faraday SOC.,46, 612 (1949).
NaClOd 1-9
Y
1 - 9
Y
+0.009 -0.079 .180 .274 ,358 .443 .523 .597 .667 .736 .784 .821
[O .626 ]
$0.009 -0.025 .056 .083
.649 .676 .706 .736 .77 .81
.107 .134 .156
.89
.202
.97
.246
1.06
.286
perchlorate, the reference value a t 4 m was taken from the same compilation. Randall and White's equation" for computing activity coefficients from vapor pressure was used to get y -h +'-h 2.3 2.3 -_ h h 2.3 -k -2
log@
=i
yi
-
&
ck
dm
h $:/' a, dm'h
(la) (lb)
Here ymis the activity coefficient a t molality (g. per 1000 g. water), m, and T~ is the same quantity at the reference molality, 1 m (or 4 m). h = 1 - 4 where qi is the osmotic coefficient. At these high concentrations the integral in equation 1 contributes an important part t o y . It was therefore (11) M. Randall and A. McL. White, J . Am. Chsm. SOC.,48, 2614 (1926).
186
M. L. MILLERAND M. DORAN
evaluated graphically from two plots, one of h/m versus m and one of h/rn'/*ver.sus m'/l BS jn equations l a and lb. The values of y so computed and of 4 are summarized in Table 11.
Vol. 60
Discussion of Results The behavior of these salts in highly concentrated solutions is an extension of their behavior a t low concentrations and shows nothing unexpected.
CONCENTRATED SALT SOLUTIONS. 11. VISCOSITY AND DENSITY OF SODIUM THIOCYANATE, SODIUM PERCHLORATE AND SODIUM IODIDE BY M. L. MILLERAND M. DORAN Contribution from the Stamford Laboratories, Research Diwision, American Cganamid Company, Stamford, Conn. Received April #Q, 1Q66
Viscosities of solutions of NaSCN, NaCIO, and NaI from low Concentrations up to saturation (or above) have been measured at 0, 30 and 50". Usin absolute rate theory, free energies, entropies and heats of activation of viscous flow have been calculated at 30". These v i u e s have been interpreted to mean that at high concentration the three salts studied take on quasi-crystalline short range local order. This interpretation is supported by computation of excess partial molal entropies.
'
4
I. Introduction Theoretical treatments of aqueous salt solutions have been confined to the dilute or moderately dilute region (not over 4 N ) . The range 6 N or above is largely an uncharted wilderness containing few experimental outposts. It seems reasonable to expect that a valid theoretical treatment, when it comes, may well stem from a theory of molten salt behavior rather than from a theory of the infinitely dilute solution. With this in mind, we should look for evidence, in the properties of very concentrated aqueous electrolyte solution of the beginning of short range local order. I n the very concentrated ranges, we should also expect to find that ion size plays a predominant role in determining solution properties. Some evidence for the existence of a t least short range local order in concentrated salt solutions has already been found in studies of (a) density,'B2 and (b) vapor pressure.* Since 1936, there have been a number of attempts to use the quasi-crystalline concept as a basis for a theory of the liquid state. (See reference 4 and references quoted therein.) Although somewhat declining in favor as applied to pure liqu i d ~ the , ~ concept of quasi-crystalline local order has much more to recommend it in the field of highly concentrated salt solutions (see Duttra6). Eyring and associates' have used, with much success, the concept of quasi-crystalline structure in correlating the properties of pure liquids. Their methods are available for use in the study of concentrated aqueous salt solutions. In this work, we have measured the viscosity of sodium thiocyanate. sodium perchlorate and sodium iodide solutions a t 0, 30 and 50" from zero (1) H. 8. Harned and B. B. Owen, "The Physical Chemistry of Electrolyte Solutions,'' 1st Ed., Reinhold Publ. Corp.. New York. N. Y.. 1943, p. 259. (2) A. F. Scott, THISJOURNAL, S6, 3379 (1931). (3) R. H. Stokes and R. A . Robinson, J . Am. Chem. Soc., 7 0 , 1870 (1948). (4) J . 9. Rowlinson and C. F. Curtiss. J . Cham. Phys., 19, 1519 (1951). (6) J. H. Hildebrand, Disc.Faraday Soc., 16, 9 (1953). (6) M. Duttra, Proc. Nall. Znsl. Sci. India, 19, 183 (1953). (7) 8. Gladstone, K. J. Laidler and H. Eyring, "The Theory of Rate Processes," McGraw-Hill Book Co., New York, N. Y., 1941, 1st Edition.
per cent. to near (or sometimes above) saturation. Although there is a large amount of viscosity data in the literature, there is very little on aqueous solutions above 5 N and even less in this concentrated range a t two or more temperatures. Among 1-1 salts only ammonium nitrate and silver nitrate, measured by Campbell and Kartzmark" and Campbell, Gray and Kartzmarkg at 25, 35 and 95' and the data on lithium chloride in "The International Critical Tables" meet the above requirements. II. Experimental Preparation of Solutions.-The source of the salts and the preparation and assay of solutions was the same as in previous work on vapor pressure.10 Measurement of Viscosity.-Viscosity measurements were made using conventional techniques. Timing was by an electric timer registering to 0.1 second. Two viscometers, an Ostwald and a Cannon-Fenske type, were used. T o guard aga.inst instrument error and to be sure that shear effects were not entering a t higher concentrations, several of the more concentrated solutions were measured in both viscometers. For a measurement, 5 =I=0.15 ml. was introduced into a viscometer and the amount weighed. The exact volume at the temperature of measurement was calculated from the density. Use of a weighed amount of solution was necessary to ensure sufficient precision. Because the volume, especially a t high concentrations, wag rarely accurately 5 ml., it was corrected to 5.00 ml. by a calibration curve prepared for each viscometer. The correction for a difference of 0.15 ml. never amounted to more than 0.8% of the flow time. Viscometers were calibrated to correct for turbulent flow and drainage using water, and 20 and 40% sucrosell solutions as standards. Sucrose solutions were prepared by weight and their concentrations checked by measurement of the index of refraction. Viscosity values for sucrose solutions were taken from the work of Bingham and Jackson.I2 The calibration equation was assumed to be of the form q / d = At B/t where t is the flow time, d is density and 7 is absolute viscosity in centipoise.
-
(8) A. N. Campbell and E. M. Kartzmark, Can. J . Chem., 30, 128 (1952). (9) A. N. Campbell, A. P. Gray and E. M. Kartzmark, ibid., 81, 617 (1955). (10) M. L. Miller and C. L. Sheridan, THIS JOVRNAL,60, 184 (1956). (11) Glycerol solutions, commonly recommended as calibrating liquids, proved too hygroscopic. (12) E. C. Bingham and R. F. Jackson, Bull. Bur. Slda., 14, 59 (1918-19).
